Many high-school chemistry texts, general science texts, and similar
authorities emphasize the distinction between “chemical” change and
“physical” change. (Often the first lesson of the school year
revolves around this distinction.) There is one major problem with
this, and a long list of lesser problems.

There is a moose on the table, i.e. a fundamental dilemma that the
authorities are apparently unable to resolve, or even to face
squarely.

In an introductory chemistry class, early in the term, it would
be nice to tell students roughly what a chemical reaction is.
That’s a reasonable goal, but many authorities go about it in a silly
way. They ask students to classify processes as “chemical” versus
“physical” based on superficial, macroscopic observations. The
books wantonly ignore the fact that chemical reactions are
ultramicroscopic ... inherently and necessarily ultramicroscopic. As
will be demonstrated below, there is not any reliable way to use
macroscopic observations to classify something as chemical, physical,
or both, or neither.

The books try to waltz around this fundamental dilemma.
This produces all sorts of nonsense, including:

First of all, the books wrongly claim that there exists a
meaningful distinction between “physical” and “chemical” change.
They point to a few extreme cases where such a distinction can be made
... while ignoring the fact that in the real world there is a vast
gray area where chemistry and physics overlap. At the microscopic
level, it is hard to draw a sharp distinction between chemistry and
physics, and at the macroscopic level it is even harder.

Secondly, nobody cares. If you did label this-or-that process as
“chemical” or “physical”, the label would have zero practical
value. This point is of course related to the previous point, because
if the label did have any practical use, that would in itself form
a sufficient basis for making the distinction.

In the course of scientific work, it is common to begin with a
few simple observations, then notice a pattern in the data, and then
formulate a rule that (a) gives a unified description of the existing
observations and (b) makes predictions about future
observations.1
So far so good.

The problem is that all too commonly, the authorities give us rules
that are grossly inconsistent with the examples they have given! Four
such rules will be exhibited below.

What’s worse is that sometimes the books ask the students to
make the generalizations on their own, generalizing from a small
number of examples. Perhaps this is supposed to give the students
practice in “thinking” and in “doing science”.

This, alas, violates one of the most basic rules of scientific
thinking. As discussed in reference 1, one should
always consider all the data, or at least a fair sampling of the
data. It is outrageous to base a rule on a handful of examples that
have been selected or contrived to support the rule, while ignoring
other examples that conflict with the rule. (See section 2 for
examples relevant to the “chemical versus physical” discussion.)

The result is that students are being taught something that is a
mockery of thinking and a travesty of science. They are learning
a bunch of baloney that will have to be unlearned later. Any good
teacher knows that unlearning something is an exceptionally difficult
and unpleasant task.

In general, any change of phase, such as
liquid ↔ vapor, is a physical change.†

In
general, any chemical reaction that converts one type molecular
entity into another type of molecular entity is a chemical change.
As a particularly simple example, consider the reaction 2F
↔ F2.

Now let us see how far we can get if we attempt to form
generalizations based on the available examples.

Attempted Rule A:
Any change initiated by simple physical or
mechanical means is a physical change.†

For instance, casual mixing is obviously
a mechanical process. Similarly cutting with scissors is a mechanical
process. Opening a valve is mechanical.

This rule runs into trouble as soon as we consider example 4.
The physical change is controlled by a change in temperature. If you
lower the temperature, the vapor condenses into liquid. If you raise
the temperature, the liquid evaporates into vapor and mixes with the
air. If you raise the temperature some more, the air/vapor mixture
ignites – or explodes – which is a chemical reaction.

From this example, and many others, we learn that sometimes
temperature initiates a physical change, and sometimes initiates
a chemical change. This already calls Attempted Rule A into
doubt.

Suppose we have an ordinary off-the-shelf bottle of carbonated
water. If you remove the cap, gas rushes out. This seems closely
analogous to example 3. But are we seeing a chemical
change, or a physical change? It turns out we are seeing
both. Before the bottle was opened, there was a four-way
equilibrium:

Some of the CO2 was in the gas phase, in the head space.

Some of the CO2 was dissolved in the water, in the form of CO2.
The dissolved CO2 was in equilibrium with the gaseous CO2.

Some of the CO2 had reacted with the water, to form
carbonic acid, i.e. H2CO3. The reaction CO2 +
H2O
↔ H2CO3 was in nontrivial equilibrium,
i.e. it did not go to completion in either direction.

Some of the carbonic acid had ionized. The reaction
H2CO3 ↔ H+ + HCO3− did not go
to completion in either direction.

Now the point is that when you reduce the pressure by removing the
cap, there are some obvious physical changes that take place, but
there are also two chemical reactions that take place. At the new
pressure, the chemicals
are no longer in equilibrium, so some of the carbonic acid
dissociates into water and CO2. Also some of the
carbonate ions convert to un-ionized carbonic acid molecules.

This is the same as example 6, but rather than merely opening
the cap and letting gas escape, we attach a piston to the top of the
vessel. This allows us to raise and lower the CO2 pressure
in the head space.

When we do this, we discover that the system is reversible. A
higher pressure of CO2 in the head space is associated with
more CO2 in solution, and hence more H2CO3 in
solution, and hence more H+ and HCO3− ions in
solution.

At this point, it should be obvious that Attempted Rule A is deader
than a doornail. Indeed, the Fundamental Tenet is also untenable.
There is absolutely no way that an incoming student in an
introductory-level course could, on the basis of casual macroscopic
observations, tell whether example 6
depends on chemical changes, physical changes, or both. (Also, as
mentioned in section 1, keep in mind that nobody cares.)

At this point, any sensible person would give up, since the
Fundamental Tenet has been discredited. But chemistry texts rush in
where angels fear to tread. Here’s another example:

Some people try to salvage Attempted Rule B by claiming that the
scissors (a) doesn’t cut any covalent bonds and therefore (b)
doesn’t “really” cut molecules but rather cuts “between”
molecules. First of all, part (a) of this claim is untrue, as you can
see from example 8 and example 9 ... and even if part
(a) were true, part (b) would violate the IUPAC definition of
molecule. This salvage attempt cannot possibly succeed.

Suppose we start with a good-sized crystal of calcite.
The crystal is beautifully clear. Split it
using a geologist’s hammer. We now have have two crystals, just
smaller. This process was carried out using an obviously
“physical” means, but the most-important step in the process
was the breaking of a huge number of chemical bonds in the
crystal. If you keep pounding on the calcite with a hammer,
you can quickly reduce it to a fine white powder.

There is no doubt that the nylon in example 8 consists of very
long macromolecules – centimeters long, sometimes many meters long,
covalently bonded from end to end – which can be readily cut by
scissors. Similarly the macroscopic calcite crystal in
example 9 is one huge covalently-bonded macromolecule. This
can be cut into lower-molecular-weight molecules by purely mechanical
means.

Minor point: This leads to unanswerable questions about
how many broken bonds is “too many”. These questions are
unanswerable because there will always be marginal cases, as you can
see from example 9. There is no practical limit as to how
finely the calcite crystal can be crushed, and therefore no practical
limit as to how many bonds can be broken. This is an example of a
“camel’s nose” argument. Once you allow a few bonds to be broken,
there is no good place to draw the line.

Actually, the situation is even worse than that. There are two or
three fundamental flaws in Attempted Rule B2. The first is a logical
flaw: When you cut a piece of nylon with scissors, the bonds that are
cut represent a small percentage of the total number of bonds ... but
the bonds that are cut represent 100% of the cutting. So if we look
at the cutting process per se, it consists entirely of chemical
changes, even though it is routinely claimed to be “obviously” a
physical change.

Another fundamental flaw is this: Not only is there no “good”
way to decide how many broken bonds is too many, there is
absolutely no place to draw the line that is consistent with the
facts. That’s because there are processes that the authorities like
to call “physical” where a majority of the chemical bonds are
broken, as we see from the following examples:

Suppose we start out with a good-sized crystal of halite, which
has empirical formula NaCl. The crystal is a single molecule,
i.e. a macromolecule, ionically bonded.
We can agree that the empirical formula
is NaCl, and the unit cell formula is NaCl, but actual
molecular formula is NaxClx, to a high degree of
approximation, for some large value of x. Values on the
order of x=1020 or even larger are commonly encountered.

Now suppose we heat the crystal so that some vapor is
formed. Under a wide range of conditions of temperature and
pressure, the vapor will contain lots of NaCl molecules, that is,
plain old Na1Cl1 molecules. This process is reversible,
in the sense that if you lower the temperature, it is possible to
regrow the crystal at the expense of the vapor.

In the NaCl molecules in vapor phase, each Na atom is bonded to
exactly one Cl
atom, and vice versa. In the crystal, each atom was bonded to
its six nearest neighbors.

This is the same as example 10, but this time we use a
covalently-bonded crystal such as silicon. The crystal is
one big molecule. Each atom is covalently bonded to
each of its four nearest neighbors.
The vapor (under suitable conditions) consists
of isolated silicon atoms – no bonds at all.

In example 10, for instance, you could say that 5/6ths of the
ionic bonds are broken. In example 11, all of the covalent bonds
are broken. These simple phase changes are widely considered to be
physical changes, yet there is no way they can be construed as
breaking only a “few” of the bonds.

You could try to defend the idea of chemical versus physical change by
saying that any bonds that get broken when something
evaporates will – arbitrarily – not be counted. However, even that
last-ditch defense leads to logical inconsistencies, as we see from
the following example, which is closely related to
example 5:

A
chemical reaction such as 2 CH3COOH ↔
(CH3COOH)2 is a chemical change.

We now invoke the principle that a process that consists of two or
three physical changes is itself a physical change. This principle
often goes without saying, but it remains true and important. This is
called the limited transitive property.

Note that I am not claiming an “unlimited” transitive property.
We know that a process that breaks a few bonds, if repeated enough
times, will break a great many bonds. This is the minor point made
above, and is not the point that I wish to emphasize.

A more interesting and more fundamental argument can be made
as follows: It turns out that
in the vapor phase, acetic acid tends to form dimers, i.e.
(CH3COOH)2 ... especially at low temperature and/or low
molar volume. This is analogous to the dimerization reaction
mentioned in example 5, but perhaps more familiar,
because it occurs under more convenient conditions of pressure
and temperature.

Now if you believe that a liquid ↔ vapor phase change
is a physical change, we have a clear violation of transitivity. We
have two clear-cut examples of supposedly physical changes that, in
combination, are equivalent to one clear-cut chemical change. The
situation is diagrammed in figure 1.

The fundamental problem here is that when people classify phase
changes as physical changes, they are implicitly assuming that the
forces that hold crystals together – and the forces hold liquids
together – are somehow different from the forces that hold molecules
together. They are assuming something that is fundamentally not true.

— — —

Believe it or not, even after all that, there are people who still try
to defend Attempted Rule B. They say it is important, because
of the following:

Attempted Supporting Rule: molecules are the “stable
particles of matter”† and therefore deserve special consideration.

Alas, that attempted supporting rule is a non-starter, as illustrated
by example 13.

Suppose we have a container of ordinary water sitting on the shelf.
We just let it sit there. Conditions of temperature, pressure,
etc. do not change.
You might think that by any reasonable definition, no physical
change is occurring, and no chemical change either.

However, note that at standard temperature and pressure, about 18 ppb
of the water is auto-ionized, i.e. dissociated into H+ and OH−
ions. (The ions are, of course, solvated.) Ions are recombining
and new ions are being formed on a very rapid timescale.

If you add heavy water (D2O) to ordinary water, you will very
quickly wind up with a lot of HDO molecules.

Water molecules are not “stable particles of matter”. Let’s be
clear about this:

At this point, you may be convinced that trying to distinguish physics
from chemistry is a pointless exercise. If so, you can stop reading
now. However, experience shows that some people are not yet
convinced.

Attempted Rule C:
A chemical process (unlike
a physical process) produces a product with different properties than
either ingredient.†

This rule is, alas, quite unreliable. There are many
counterexamples, including example 9. The powdered calcite
has several properties not shared by the original large crystal. For
starters, it is a different color. Other counterexamples include
example 14, example 15, and example 16.

Suppose we take some copper and alloy it with a little bit of tin.
Most people consider this a purely physical mixture. That is to
say, there is no sign of any intermetallic compound being formed.
However, the product – called bronze – has properties are definitely
not what you would predict just by averaging the ingredients:

It has a lower melting point than either ingredient.

It has a different color.

It has different thermal-expansion properties.

It has greater hardness than either ingredient.

It has wildly different electrical conductivity, especially
at low temperatures.

As a more prosaic example, consider what happens if you mix the
yellow dye from an ink-jet printer with the cyan dye. The mixing
process is not a chemical process; basically it is just a mixture,
rather like example 1.

However, the mixture has a different color than either ingredient:
it’s green.

Perhaps an even better
example is the following:
Start by asking students what what was the color of Napoleon’s
white horse. This is a silly question ... a proverbially silly
question. The answer, of course, is white.
Next, ask them what color is
the stuff that comes in an ordinary bottle of yellow food coloring.
Many of them will assume that it is yellow. A few of them may know, based
on experience, that the right answer is otherwise. In fact,
the stuff in
the bottle is red. Really, deeply red. If you dilute
it enough, it becomes orange, and if you dilute it even more,
it becomes yellow.

If you do the experiment, many students will tell you that diluting
the yellow food coloring must involve some sort of “chemical
change” because they’ve been taught that color change implies
chemical change. That is, alas, dead wrong in this case.

To understand what is really going on, dilute the yellow food
coloring to a moderate degree and put it in a white
bowl with sloping sides. You will observe that in thin layers, the
liquid is perceived as yellow, while in moderately thick layers it is
perceived as orange, and in yet thicker layers it is perceived as red.

Remark #1: It should be obvious that there is no “chemical
change” involved in going from a thin layer to a thick layer.

Remark #2: There are good reasons – excellent physical reasons – why
the thin layer should have a different perceived color from the
thick layer. The physics of light absorption is nonlinear. See
reference 2.

Saving the strangest for last, we have this gem. Some people
advocate the following rule:

This rule is supported by example 1, but contradicted by a host
of other examples. For example, the chemical reaction mentioned in
example 5, namely 2F ↔ F2, is commonly
carried out under conditions where it is reversible for all practical
purposes. Raising the temperature and/or raising the molar volume
shifts the equilibrium to the left, whereas lowering the temperature
and/or lowering the molar volume shifts the equilibrium to the right.
Note that this provides another disproof of Attempted Rule A, since
we have a chemical reaction that can be initiated – in either
direction – by a simple, physical, mechanical change, i.e. changing
the volume.

It may be that the first few reactions that you saw in high-school
chemistry class were irreversible, but in the real world, many
reactions proceed both forwards and backwards. Example 7 is
is another illustration; example 17 is yet another.

Consider the electrochemical reactions in a storage battery.
By the physical process of turning the crank on a dynamo, you
can run the reactions forward. By turning the crank the other
way, you can run the reactions backward.

The other half of Attempted Rule D is wrong, too; there are plenty
of so-called “physical” processes that are, in practice,
irreversible. Crushing a calcite crystal to powder
(example 9) is irreversible; there is no easy way to un-crush
it. Similarly, cutting paper with scissors (example 2) is
irreversible; there is no way to use the scissors “backwards” so as
to un-cut the paper. More importantly, there are plenty of
truly physical processes, involving no chemical changes
whatsoever, that are irreversible. Stirring a liquid is
thermodynamically irreversible. An ideal gas rushing through an
orifice from a high-pressure region to a low-pressure region is
thermodynamically irreversible. Heat flow along a bar that is hot at
one end and cold at the other end is thermodynamically irreversible.
Other examples abound.

The preceding discussion should make it clear that there is no useful
distinction between chemical change and physical change ... except
possibly in a few extreme cases.

Any students who are dissatisfied with the usual “textbook”
discussion of this topic should be congratulated. It shows they have
some critical-thinking ability.

Too much of intro-level chemistry is devoted to rote learning of
notions that are not really correct, and would be worthless even if
they were correct.

We should concentrate on meaningful rather than meaningless
classifications.

++ I am perfectly willing to study the geology of eastern Colorado,
and to classify the rocks according to their observable properties.

++ I am perfectly willing to study the geology of western Kansas,
and to classify the rocks according to their observable properties.

- - I am not willing to force students to pretend there is a
distinction between the geology of eastern Colorado and the
geology of western Kansas.

Let me explain the meaning of this metaphor:

There is a sharp,
legal line between Kansas and Colorado. They do not overlap. The
line is, however, arbitrary. It has no effect on the geological
“ground truth”.

There is tremendous overlap between “physics” and
“chemistry”. There is no sharp dividing line. Even if somebody managed
to lay down such a line, it would be completely arbitrary. It would
have no effect on the ground truth. The atoms and molecules are what
they are, and they do what they do. They do not recognize any
distinction between chemistry and physics.

In some jurisdictions, teachers are required to cover the topic
“chemical versus physical change”. Sometimes this requirement is
direct, and sometimes it is indirect, in the sense that the topic is
included on the mandatory standardized tests.

I don’t recommend ducking the issue. We need to teach students how to
deal with bad ideas that are likely to show up on the test.
I take the direct approach. I tell them:

If it’s not worth doing, it’s not worth doing right.

Don’t pollute your brain.

Here’s what I mean by that: Ordinarily, the right way to learn
things is to take each new idea and turn it over in your mind, to
see how the new idea is connected to everything else you know.
However, when faced with a pseudo-idea that doesn’t make sense, you
should not try to make sense of it. Just learn the pseudo-idea by
rote. Don’t think about it. If you think about it, you will just
pollute your brain.

Also: ordinarily cramming is a bad idea. Anything you “learn”
one hour before the test you will forget one hour after the
test. However, for pseudo-ideas, that’s just what you want.
There are only a few pseudo-ideas that show up on the test,
and I can tell you what they are, so make a list. Cram them
into rote memory right before the test, and forget them right
afterward. The test cannot possibly ask any deep questions
about these pseudo-ideas, so rote memory will be quite sufficient.

Meanwhile we should agitate to get the nonsense removed from the
tests.

Also: Remember that elections have consequences. Find out about the
candidates, even in off-year elections, and even in down-ballot races
such as school board. Then be sure to vote. Otherwise the wackos will
get elected.

Please let’s stop foisting ideas of physical versus chemical change
onto the students. Those ideas are worse than worthless. The less
said about them, the better.

The only reason why anyone would reasonably be interested in such
ideas would be if they lived in the 19th century, before there was any
useful understanding of atoms.

The 19th century has been over for a while now. Wake up and smell the
atoms. Anything you could possibly explain in terms of a
physics-versus-chemistry distinction can be explained infinitely more
clearly by saying what the atoms and molecules are doing.
Especially in an introductory course, students should see the best
evidence, not the most ancient evidence. See reference 3.
A useful valid microscopic story is preferable to a useless
invalid macroscopic story.

There is a saying, “You can’t beat something with nothing”.
The question therefore arises, if we aren’t going to spend the
first week of class talking about “chemical” changes versus
“physical” changes, what should we talk about instead?

The answer is simple: Talk about real things. Talk about protons,
neutrons, and electrons, and how they form atoms
(reference 4). Talk about atoms and bonds and molecules.
See reference 5 for suggestions on how to introduce such ideas
at the high-school or even pre-high-school level. Talk about energy
and entropy (reference 6). Talk about dimensional
analysis and scaling laws (reference 7 and
reference 8). Talk about things that really matter.